Towards Magic Photoacids â Proton Transfer in Concentrated Sulfuric

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A: Kinetics, Dynamics, Photochemistry, and Excited States

Towards Magic Photoacids – Proton Transfer in Concentrated Sulfuric Acid Daniel Maus, Alexander Grandjean, and Gregor Jung J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.8b09974 • Publication Date (Web): 25 Oct 2018 Downloaded from http://pubs.acs.org on October 26, 2018

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The Journal of Physical Chemistry

Towards Magic Photoacids Proton Transfer in Concentrated Sulfuric Acid Daniel Maus, Alexander Grandjean and Gregor Jung* Biophysical Chemistry, Saarland University, Campus B2 2, 66123 Saarbruecken, Germany

ABSTRACT

Photoacids are the most convenient way to deliver protons on demand. So far, their photoacidity allows for studying excited-state proton transfer (ESPT) only to protic or strongly basic solvent molecules. The strongest super-photoacids known so far exhibit excited-state lifetimes of their conjugate base in the order of 100 ps before recapturing the proton again. Here, we describe how we developed a new aminopyrene based super-photoacid with an excited-state lifetime of its conjugate base of several nanoseconds. It will be shown by fluorescence titration and via Förster cycle that the excited-state acidity is as high as concentrated sulfuric acid and thus exceeding any previous photoacidity by several orders of magnitude. Its outstanding chemical stability and fluorescent properties make it suitable for time-resolved proton-transfer studies in concentrated mineral acids and organic solvents of low basicity.

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INTRODUCTION Theodor Förster already interpreted the pH-sensitive dual-emissive behavior of some aromatic fluorophores as photoacidity, which is an increase in acidity by several orders of magnitude after electronic excitation1,2. Time-resolved studies on proton transfer exploit short laser pulses to trigger the release of a proton from these so-called photoacids and the kinetics may be subsequently probed by fluorescence spectroscopy.3–9 The excited-state acidity (pKa*) of photoacids is the key parameter for such an experiment as it determines which molecules can act as proton acceptor. For example, the use of pyranine (HPTS) as photoacid is limited to water-based systems as its pKa drops from 7.3 to only 1.3 upon excitation.3,10 By means of super-photoacids (pKa* < 0) excitedstate proton transfer (ESPT) also takes place in some organic solvents like alcohols and DMSO.8,9,11,12 It has been shown in naphthols and pyrenols that the acidity of photoacids can be increased by introducing electron-withdrawing substituents.11,13–15 As a result of lowering the electron-density of the acidic group, the super-photoacids 5,8-dicyano-2-naphthol (pKa* = −4.5)14 and tris(1,1,1,3,3,3-hexafluoropropan-2-yl)-8-hydroxypyrene-1,3,6-trisulfonate (pKa* = −3.9)11 have been developed. Introducing positive charges into a photoacid causes a stabilization of the negatively charged conjugate base leading to even stronger photoacids like N-methyl-6hydroxyquinolinium (NM6HQ+)16–18 with a pKa* of −4.0 as well as the quinone cyanines QCy7 (pKa* = −5.7)19 and QCy9 (pKa* = −8.5)20,21, which are the strongest published super-photoacids so far. For an up-to-date overview of the topic of photoacidity see for instance a recent review22 and a special issue23. However, in order to extend the proton transfer studies to other solvents (e.g. acetone and acetonitrile) or mineral acids, stronger photoacids are needed. Therefore, reaching pKa* ≈ −10 can be considered as a milestone. This value corresponds to the Hammett acidity function24–26 of concentrated H2SO4 (96 %, H0 ≈ −10)26, which already serves as a benchmark in


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ground-state acidities:27,28 Stronger ground-state acids, including the well-known magic acid, are called super-acids.27 Another important parameter is a sufficiently long excited-state lifetime of the conjugate base in order to allow complete diffusional separation into the free ions. However, this parameter strongly depends on the solvent and may play a minor role, for example, in neat water. While the excited conjugate base of QCy9 has a relatively short fluorescence decay time of about 90 ps, the high ESPT rate (kESPT ≈ 1 ∙ 1013 s-1) of the photoacid provides the fully separated ions in water within just a few picoseconds.21 Nevertheless, it has been shown in the past that once mixtures of mineral acids or organic solvents with water are used instead of neat water, the rate constant for ESPT significantly reduces with a decrease in the amount of water molecules.21,29–31 Interestingly, a deviation from this behavior has been described recently. The photoacid 2-naphthol-8-sulfonate showed a higher ESPT rate in acetonitrile-rich CH3CN/H2O mixtures than in neat acetonitrile. The authors explained this observation by the presence of a water bridge between the 2-OH and the basic 8-sulfonate, enhancing the ESPT process.31 However, we assume that in concentrated sulfuric acid and dry, low basic organic solvents the ESPT rate becomes rather small due to the absence of free water as acceptor. Additionally, in concentrated H2SO4, we assume a large amount of protons available to enhance the reprotonation reaction in the excited state, which eventually increases the importance of a longer fluorescence lifetime of the conjugated base. Pyrene-derived photoacids perfectly match the requirements for the development of new superphotoacids, because of their variable derivatization, high chemical stability and long fluorescence lifetimes around 5–6 ns.11,32 Here we modified the less noticed, but highly photoacidic33,34 aminopyrenes by electron-withdrawing side groups. In these systems, the photoacidic species (−NH3+*) must first be generated by protonation in the ground state, before ESPT can occur to


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form an amino species (−NH2*), its conjugate base.2,5,33 It will be shown by optical titration and via Förster cycle that pKa* values close to −10 are obtained, thus exceeding any previous photoacidity by several orders of magnitude. METHODS General. All chemicals were obtained from commercial suppliers and were used without further purification. Absorption spectra were recorded with a double beam Jasco V-650 spectrophotometer. A Jasco FP-6500 spectrofluorometer was used for recording the fluorescence emission spectra. The fluorescence quantum yields were measured in scan mode (5 nm increments around the absorption maximum) on a Hamamatsu Absolute PL Quantum Yield Spectrometer C11347. All quantum yields were averaged from the values at the absorption maximum and the two above and below it (five values total). All samples for fluorescence related measurements were prepared with an optical density of about 0.1 at the absorption maximum of the dye unless otherwise noted. Synthesis. The two step synthesis of the new photoacid tris(2,2,2- trifluoroethyl)-8-aminopyrene1,3,6-trisulfonate (2) and its N-derivative 8-aminopyrene-N, N, N’, N’, N’’, N’’-hexamethyl1,3,6-trisulfonamide (3) (Scheme 1) followed a modified combination of two published routines.11,35


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Scheme 1. Synthesis of the Aminopyrene Derivatives 2 and 3.

1-Aminopyrene (1) (>98 %, TCI Germany) in chlorosulfonic acid provided at first the nonfluorescent trisulfonyl chloride. The isolated dark red solid was then converted via a nucleophilic substitution with either 2,2,2-trifluoroethanol or dimethylamine into either 2 or 3. The compounds were obtained as orange solids in 55 % (on gram scale) (2) and 61 % (3) overall yield. The products were identified by 1H-, 13C- and 19F-NMR (for 2) spectroscopy and mass spectrometry. Further details of the synthesis and the characterization can be found in the Supporting Information. Absorption Titration. According to the literature24,36, H0 / pH -values were adjusted between −5.02 and 0.39 by dilution of perchloric acid (70 %, Acros Organics) with water. Commercial buffers (Carl Roth) were used for pH = 2 (Citric acid / HCl / NaCl), pH = 3 and pH = 4 (Citric acid / NaOH / NaCl), pH = 5 and pH = 6 (Citric acid / NaOH). A small aliquot of a stock solution of 2 or 3 in the concentrated acid was added to each solution. The amount of aliquot added was taken into account when preparing the specific H0 / pH-values. Fluorescence Titration. According to the literature25,26, H0-values were adjusted from −12.62 to −6.10 by either dilution of concentrated sulfuric acid (96 %, Carl Roth) with water or by adding commercial oleum (20−30 % free SO3, Acros Organics) to concentrated sulfuric acid (96 %, Carl


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Roth). Small aliquots of a stock solution of 2 in concentrated sulfuric acid were added. Those amounts were taken into account when preparing the specific H0-values. Time-correlated single-photon counting (TCSPC). As light sources for the TCSPC measurement were used a pulsed laser diode (LDH-D-C-375, PicoQuant, λexc = 375 nm, 20 MHz, pulse width < 100 ps) in combination with a laser diode controller (PDL 800-D, PicoQuant) and another pulsed laser diode (LDH-P-C-470, PicoQuant, λexc = 470 nm, 20 MHz, pulse width < 100 ps) in combination with a laser diode controller (PDL 808 MC SEPIA, PicoQuant). A photon counting detector (PDM series, Micro Photon Devices) and a photon counting device (PicoHarp 300, PicoQuant) were used for detection. The collected data was analyzed using the SymPhoTime software (PicoQuant). A diluted colloidal silica solution (LUDOX TM-50, SigmaAldrich) and an aqueous solution of erythrosine B (quenched with KI) was used for recording the instrumental response function (IRF), which had an overall FWHM ≈ 300 ps (λdet = 417−477 nm) and FWHM ≈ 100 ps (λdet = 540−600 nm). RESULTS AND DISCUSSION The ground-state pKa of compound 2 and 3 (see SI) were derived from an absorption titration. The Hammett acidity function was used as extension to the pH-scale in which the H0 values can be adjusted by a dilution series of either sulfuric or perchloric acid.24–26 The latter one was used here because of its lower viscosity and heat development by dilution with water.37,38 Figure 1a shows the absorption spectra of 2 including the H0 / pH dependent interconversion of the two species with two isosbestic points at λ = 333 nm and λ = 398 nm.


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Figure 1. Raw absorption titration spectra of 2 in perchloric acid / buffers (a) and the resulting titration curve (b) of −NH2 (green dots, dotted curve) and −NH3+ (purple dots, dashed curve) according to equation (2a) and (2b). As one would expect, the NH3+-species, indicated by the structured band around λmax = 379 nm, is more likely present at higher acidities. The broad absorbance of the NH2-band (λmax ≈ 487 nm) appears only at comparatively lower acidities. The absorption band of the amine shows a small variation of its maximum, which might be caused by the use of different buffer systems and the


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tendency to form aggregates at lower acidities (pH > 0). The absorbance values A at λmax of the two species were taken at each H0 / pH from the raw spectra to calculate the molar fraction ratio R for each species separately via equation (1). Amin and Amax represent the minimum and maximum absorbance obtained in the titration. 𝐴 ― 𝐴𝑚𝑖𝑛

(1)

𝑅 = 𝐴𝑚𝑎𝑥 ― 𝐴𝑚𝑖𝑛

These R-values were plotted (Figure 1b) against the H0 / pH and fitted by the dashed / dotted curves according to equation (2a) and (2b). For the sake of simplicity, only the pH appears in the two equations, which, however, has to be replaced by H0 in the high acidic range. 𝑅(NH3+ ) =

1 pH ― p𝐾a

1 + 10

𝑅(NH2) = 1 ―

1 pH ― p𝐾a

1 + 10

(2a)

(2b)

The molar fraction ratio curves intersect at their inflection point at R = 0.5 giving a pKa value of −0.5±0.1 for 2. Thus, in the ground state 2 is about 105 times more acidic than its hydroxyl analogue (pKa* = 4.7)11, confirming the same tendency as observed within the pair of phenol (pKa = 9.9) and anilinium (pKa = 4.6).39 In any case, excitation of −NH3+ of 2 at pH / H0 > −7.75 (Figure 2a) leads to a strong yellowish fluorescence emission at λem = 547 nm (Figure 2b).


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Figure 2. Normalized absorption (a) and emission spectra (λexc = 350 nm) (b) of 2 in hydrochloric acid (black curve), aqueous perchloric acid (red curve) and perchloric acid in glacial acetic acid (blue curve). Only in HClO4 (70 %, H0 = −7.75)36 a faint bluish emission at λem = 399 nm was noticed for 2. Interestingly, the protonated sulfonamide derivative 3 (pKa = 1.2±0.1, see SI) lacks any significant fluorescence and therefore does not serve as photoacid, although its conjugate base shows a bright yellowish emission as well. The reason of this is still subject of further investigations.


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Several methods exist for the determination of pKa*.11,19,40 The most convenient and most often used way is the calculation of the pKa-drop (ΔpKa) via the Förster cycle (noted as pKa*(Förster)) (equation (2)).2,11,19,20

Δp𝐾a =

(ℎ𝜈NH2 ― ℎ𝜈NH + ) 3

𝑘𝑇 ln(10)

(2)

Here, ΔpKa depends on the energy difference of the 0-0-transitions of the photoacid and its conjugate base. From the experimental wavelengths of the absorption (λNH3+ = 391 nm, λNH2 = 487 nm) and fluorescence emission (λNH3+ = 399 nm, λNH2 = 547 nm) spectra of 2 in aqueous perchloric acid, values of −10.5 (absorption) and −14.2 (emission) were calculated, which on average give a change in the acidity of ΔpKa = −12.4. The use of the experimental ground state acidity resulted in a pKa*(Förster) of −12.9, thus exceeding the similarly determined excited-state acidity of QCy920 by more than four orders of magnitude. The experimental verification of this value requires a fluorescence titration (Figure 3a) in an acid of this strength. Therefore, in order to shift the protolytic equilibrium in the excited state completely to the side of the photoacid, at least concentrated sulfuric acid (96 %) is needed. Although perchloric acid is known to be significantly more acidic than sulfuric acid24,36,41, concentrations higher than the commercially available 70 % are known to be very hazardous because of its increased oxidizing potential and spontaneous explosive behavior.42 In addition to the increased safety, the major advantage of sulfuric acid is the already known broad range of adjustable Hammett acidities in the super-acidic range and the easy access to anhydrous sulfuric acid via addition of fuming sulfuric acid (oleum).24–26 In any case, one is leaving the region where pH-considerations for aqueous solutions do apply.41 Starting from commercial concentrated


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sulfuric acid (96 %, fluorescence emission ratio INH3+/INH2 ≈ 1:1), dilution of the acid by water shifts the protonation equilibrium in the excited state towards the amine (λmax = 546 nm).

Figure 3. Raw fluorescence emission titration spectra (λexc = 350 nm) of 2 in concentrated sulfuric acid (a) and the resulting titration curve of −NH2 (green dots, dotted curve) and −NH3+ (purple and red dots, dashed curve) (b). The error bars were calculated from the uncertainty of the amount of free SO3 in commercial Oleum (horizontal) and the standard deviation from three independent titrations (vertical).


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Whereas addition of oleum to concentrated sulfuric acid shifts the emission ratio to the side of the NH3+-species (λmax = 395 nm), the fluorescence of the conjugate base of compound 2 disappears at about 100 % sulfuric acid (H0 ≈ −11.1). In addition, a decrease in the −NH3+ fluorescence intensity can be observed by further rising acidity of the medium. Using Oleum for H0 < −10.2, the color of the sample solutions began to appear brownish (presumably due to the presence of free SO3 and decomposition of marginal organic impurities), which may result inner filter effects being involved. The fluorescence quantum yield Фfl = 0.56 (−NH3+), recorded in 100 % sulfuric acid, therefore appears as a lower limit of the true value. In contrast, the fluorescence quantum yield for −NH2 is Фfl = 0.91, which is equal to 2 in various organic solvents (data not shown). It should be noted that decomposition by hydrolysis of 2 in solutions of H2SO4 (100 % and below) is negligible on a time-scale of days (see SI, Figure S20). Analogous to the titration of the ground state, the molar fraction ratios were plotted vs H0 after calculation via equation (1) with A as fluorescence emission intensities at λmax (Figure 3b). The curves were fitted by equation (2a) and (2b), respectively, in which the values of NH3+ at the highest acidities (red dots) were masked. It is noticeable that the intersection of the curves is located slightly above R = 0.5 at a molar fraction ratio of R ≈ 0.65, possibly due to the abovementioned inner filter effect. Furthermore, the measured R-values at H0 > −10 show some kind of “flattened” titration curves, indicating a higher presence of the photoacid and a lower presence of the conjugate base than expected. Both effects may caused by the sharply varying occurrence of different species in concentrated sulfuric acid and the associated highly significant change in viscosity and ionic strength of suchlike solutions.37,43 However, this behavior is still subject of further investigations. Because of the deviation of the intersection from R = 0.5 two different pKa* values can be determined from the inflection points: pKa* = −9.6±0.1 (from −NH3+) and


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pKa* = −10.2±0.1 (from −NH2), which on average give an excited-state acidity of pKa* = −9.9±0.1. Both, pKa*(Förster) and pKa* indicate a very high photoacidity of 2, but differ significantly. One aspect that should not be ignored is that the determination of pKa*(Förster) does not take into account any change of molecular geometry and solvation relaxation in the photochemical cycle.11,14,44 For verification of ESPT in concentrated sulfuric acid, the fluorescence of 2 in concentrated sulfuric acid and its deuterated analogue (Figure 4a) was studied in more detail by time-correlated single- photon counting (TCSPC) (Table 1). Table 1. Fluorescence Lifetimes of 2 in H2SO4 and D2SO4 H2SO4

D2SO4

λdet / nm

τ / ns

arel

τ / ns

arel

417−477

1.4±0.3

1

2.8±0.3

1

1.4±0.1

−0.99±0.01

2.7±0.1

−1±0.01

5.3±0.1

1±0.01

5.4±0.1

0.95±0.01

540−600


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Figure 4. Steady-state fluorescence spectra (λexc = 350 nm) (a), photograph (b) and TCSPChistograms (c,d) of 2 in H2SO4 (green curves) and D2SO4 (blue curves): λexc = 375 nm, λdet = 417−477 nm (c); λexc = 375 nm, λdet = 540−600 nm (d). The fluorescence decay of the photoacid itself (Figure 4c, λdet = 417−477 nm) could be monoexponentially fitted with a lifetime of τ = 1.4±0.3 ns in H2SO4 and τ = 2.8±0.3 ns in D2SO4. The TCSPC-histogram of the NH2-species after excitation of the photoacid (Figure 4d, λdet = 540−600 nm) shows the highly significant proving two exponential behavior with a rising


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and a decaying component. The lifetime of the conjugate base of 2 in organic solvents (data not shown) and other pyrene-derivatives of similar constitution11 agree well with the longer decay times τ = 5.3±0.1 ns (in H2SO4) and τ = 5.4±0.1 ns (in D2SO4). The shorter time constants τ = 1.4±0.1 ns (in H2SO4) and τ = 2.7±0.1 ns (in D2SO4) with negative amplitudes (Table 1, arel) represent the population kinetics of the excited amine by ESPT. Here it can be seen that the fluorescence lifetime of several nanosecond for the conjugate base has proven to be very useful in order to allow the full diffusional separation into the free ions in concentrated sulfuric acid. The kinetic isotope effect (KIE) for ESPT of about 1.9±0.2 is in agreement with other systems.19,20 Actually, the KIE can clearly be seen with the naked eye by the fluorescence color of the solutions (Figure 4b). In a separate titration experiment (see SI, Figure S3), we verified that the actual proton acceptor indeed is HSO4-, formed from the dissociation reaction with residual water according to equation (3).43 H2SO4 + H2O ⇄ HSO4- + H3O+

(3)

However, there are still unanswered questions about the mechanisms of proton transfer in anhydrous mineral acids, such as the flattened fluorescence titration curve of 2 in concentrated sulfuric acid. CONCLUSIONS In summary, excited compound 2 (−NH3+) shows an acidity as high as concentrated sulfuric acid, which makes it up to five orders of magnitude stronger than its parent compound APTS (calculated pKa* ≈ −7)33,34 and thus the strongest photoacid reported so far with accompanying long observation window. Both pKa and pKa* have been experimentally determined by optical titration experiments. Time-resolved spectroscopy in concentrated mineral acids proves ESPT to


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acceptors with pKa / H0-values ≥ −10. In contrast to the more common photoacids, in which the photoacidic species (R-OH) can also be available in neutral solutions, the photoacidic species of compound 2 must first be generated by protonation in its ground state. However, 2 is a suitable tool for time-resolved proton transfer experiments in concentrated mineral acids. Besides that, we currently establish the conditions for water-free spectroscopy to follow proton transfers to molecules of low basicity and proton-catalyzed reactions. We also will address the question whether such strong photoacids show an inverted region for the kinetics of proton transfer, in the sense of its description by Marcus-theory.45–48 ASSOCIATED CONTENT Supporting Information. The following files are available free of charge. Additional information on the dye synthesis and characterization, titration experiments, chemical stability experiment, NMR and mass spectra (PDF) AUTHOR INFORMATION Corresponding Author *Gregor Jung. E-mail: [email protected] Notes The authors declare no competing financial interest. ACKNOWLEDGMENTS


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The German Research Foundation (DFG JU650/7-1 and DFG JU650/8-1) financially supported this work. We like to thank Reiner Wintringer, Service Center Mass Spectrometry, Saarland University, for recording the mass spectra. We also like to thank PicoQuant GmbH for the free provision of a LDH-D-C-375 diode laser and a PDL 800-D laser driver for this publication. REFERENCES (1)

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